Warm Up

Draw the Bohr Model for Aluminum and Neon.

Electrons in Atoms

Unit 2- Continued

Everything you ever wanted to know about where the electrons hang out!

Section 1: Early 1900’s

Scientists started doing a lot of experiments looking at the absorption and emission of light by matter.

Found that there is a relationship between light and an atom’s electrons.

Light behaves as a wave

Transfer of energy

Draw the Wave!

• Amplitude: height of the wave from the origin to the crest• Wavelength ( ) : the distance between the crests (m, cm,

nm)• Frequency (v): number of waves to pass a given point per

unit of time (waves/second = Hz)

An Important Relationship

The frequency and wavelength of all waves, including light, are inversely related.

As the wavelength of light increases, the frequency decreases.

Electromagnetic Radiation

• Includes radio waves, radar, microwaves, visible light, infrared light, ultraviolet light, X-rays, and gamma rays

Wave Particle Duality

http://www.youtube.com/watch?v=DfPeprQ7oGc

The Photon

Photon- a particle of electromagnetic radiation having no mass, carrying a quantum of energy.

Photoelectric Effect

Looks at the emission of electrons from a metal when light shines on the metal.

Light causes electrons to be ejected from the metal.

So, what happens when photons hit an atom and eject an electron?

The electron goes from the ground state to an excited state.

As the electron returns to the ground state, it gives off the energy that it gained- LIGHT

Energy Levels

• Energy levels are

not evenly spaced• Energy levels

become more closely

spaced the greater

the distance from the

nucleus

Another Cool Illustration

Color

The energy given off has a certain wavelength.

Wavelength determines the colors that we see.

Flame Test

Look at the color produced by the flame…

Determine the wavelength by comparing the color to those in the visible range on the Electromagnetic Spectrum.

Warm Up

You have two different samples… sample A. glows red and sample B. glows violet.

a. Draw what the waves might look like?

b. Which has the longer wavelength?

b. Which has the smaller frequency.

Atomic Spectra

White light is a combination of all the wavelengths in the visible range of the Electromagnetic Spectrum.

Spectral Tubes

Each element has a unique line-emission spectra

Atomic Line Spectrum

Interpretation of Atomic Spectra• The line spectrum is related to energy transitions in

the atom.Absorption = atom gaining energyEmission = atom releasing energy

• All samples of an element give the exact same pattern of lines.

• Every atom of that element must have certain, identical energy states

Atomic Spectrum Activity

Using Atomic Spectral DataBohr Model

• Electrons orbit around a nucleus

• Each orbit has a fixed energy and because of this cannot lose energy and fall into the nucleus

• Energy Level of an electron is the region around the nucleus where the electron is likely to be moving

This helped explain the spectral lines

Absorption- the electron gains energy and moves to a higher energy level.

Emission- when the electron falls to a lower energy level.

Schrodinger Wave Equation

Developed an equation that treated electrons as waves and described the location of electrons.

Helped lay the foundation for modern quantum theory (atomic model).

The Quantum ModelFinally– the truth (as we know it!)

• Electrons can behave as both waves and particles.

• Electrons can be considered waves with specific frequencies confined to the space around the nucleus.

• Electrons can also be considered negatively charged particles.

Quantum Theory

• Estimates the probability of finding an electron in a certain position

• We denote the position of the electron as a “fuzzy” cloud

• This volume of space where an electron is most likely to be found is called an orbital.

• The atomic orbitals have distinct shapes

Work on Wave WS- 15 min

Go to shape ppt